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Author: Subject: Small scale production of H2SO4 in the amateur lab
macckone
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[*] posted on 13-9-2020 at 10:03


Quote: Originally posted by clearly_not_atara  
If I were to guess, this is made possible because the ionic reactions are like this:

PbCl2 (Ksp 1.7e-5) + CaSO4 (Ksp 4.9e-5) >> PbSO4 (Ksp 2.1e-8) + Ca2+ (aq) + 2 Cl- (aq)

PbSO4 + HCl + Cl- >> PbCl2 + HSO4- (aq)

That is, the sulfate ions are protonated in the acidic hydrochloric acid solution, but not in the neutral solution. The success of the calcium-to-lead metathesis thus depends on a neutral pH, while the lead sulfate-to-chloride metathesis happens at acidic pH, where the sulfate ion is protonated to bisulfate (pKa ~2).


This changes with ion concentrations and removal as well.
Both reactions are equilibrium, so if you can remove the calcium chloride, the reaction proceeds reasonably well.
If you remove the sulfuric acid the second reaction proceeds as well.
Hence my thought on the soxhlet extractor where you have continuous if slow removal of reactants from one side of the equation.
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AJKOER
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[*] posted on 24-9-2020 at 08:02


Came across an interesting electrochemical based reaction to quote:

"2 FeCl2 + 2 H2O + 1/2 O2 --> Fe2O3 + 4 HCl (1)

Hydrochloric acid is regenerated, and commodity oxides of iron are obtained as a result of the reaction [8, 9]."

Source: http://rudmet.net/media/articles/Article_CIS_vol.15_18_pp.28...

Turning to sulfate chemistry, with respect to ferric sulfate by Atomistry at http://iron.atomistry.com/ferric_sulphate.html to quote:

"Upon dilution, ferric sulphate solutions readily undergo hydrolysis, precipitates being obtained which, however, have no well-defined composition.
A study of the electric conductivities of aqueous solutions of the salt indicates that the hydrolysis proceeds in two stages, embodying (1) a rapid change unaccompanied by precipitation, and (2) a slower change, progressing at a measurable rate, and accompanied by the production of a so-called basic salt. Colloidal ferric hydroxide does not appear to be formed during hydrolysis, the salt thus differing from ferric chloride and nitrate...In dilute solution ferric sulphate is reduced by metallic iron to ferrous sulphate."

More recent research on the action of O2 on iron salt solutions, for example, this fully free available 2008 reference: "Air Oxidation of Ferrous Iron in Water" by Ahmet Alıcılar, et al, at http://www.jieas.com/fvolumes/vol081-5/3-5-11.pdf.

“Abstract: Air oxidation of ferrous iron in water was studied...Thereafter, the experiment was successively repeated by blowing air to the solution without and with inert packing. Lastly, the catalytic effect of ferric hydroxide was investigated. While the maximum yield of 86 % is catalytically achieved by blowing air at a neutral medium, the oxidation was almost completed in an alkaline solution even at stationary oxidation was almost completed in an alkaline solution even at stationary atmosphere. The reaction was first order with respect to Fe2+. “

And further:

“Oxidation of iron is achieved by addition of chemical oxidants. However, it can be easily and low cost carried out by contact with air (Wong, 1984). During oxidation of Fe2+ salt aqueous solutions, poor soluble compounds including Fe3+ oxides are formed (Domingo et al., 1994). The composition of precipitate formed depends on numerous parameters such as temperature, pH, concentration, feed rate and anion nature (Das & Anand, 1995; Tolchev et al., 2002)....The oxidation kinetics of Fe(II)(aq) species has been previously reviewed by many workers (Wehrli, 1990; Zhang et al., 1992). The stoichiometry for the overall oxidation of Fe2+ ions by O2 is given by Eq. (1) (Burke & Banwart, 2002).

O2(aq) + 4Fe2+ + 6H2O ↔ 4FeOOH(s) + 8H+ (1) “

So, similar to FeCl2, the slow air pump oxidation of a RT dilute aqueous FeSO4 and metal Iron mix to the point of possible dilute H2SO4 creation may be a subject of my future investigations. Substituting (NH4)2SO4(aq)/Fe for FeSO4(aq)/Fe, may also be interest.

[Edited on 24-9-2020 by AJKOER]
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Σldritch
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[*] posted on 30-9-2020 at 00:33


I just had a crazy idea of how to make sulfuric acid. As we all know sulfuric acid can be made by adding oxalic acid to Ferrous Sulfate. Ferrous Sulfate being essentially free oxalic acid is what makes this an expensive process. But what if you made the oxalic acid in situ? If you put out a bowl of Ferrous Sulfate, sucrose and a catalytic amount of nitric acid in the dark one might expect the following reactions to take place:

C12H22O11 + 12 HNO3 = 6 C2O4H2 + 11 H2O + 12 NO
NO + O2 + H2O = HNO3 (catalyzed by iron in solution)
C2O4H2 + FeSO4 + 2 H2O = FeC2O4·2H2O + H2SO4

With the overall reaction being:
C12H22O11 + 9 O2 + 6 FeSO4*5H2O = 6 FeC2O4·2H2O + 6 H2SO4 + 23 H2O

The main problem may be overoxidation of the sugar and the slow reaction between sugar and nitric acid. The former is catalyzed by light/iron couple which should be easy to avoid unless iron by itself can do it too. Oxalate may not be rremoved fast enough due to solubilizing effects from the sugar. The latter problem could probably be solved by adding more nitric acid. Nitrogen oxides may be lost but conditions should be very favorable for reoxidation so it may be manegable.

Just like with the oxalic acid process a nice bonus is the Ferrous Oxalate dihydrate byproduct which may be useful to make iron or iron oxide powder though it may be diffucult to contain it 100g+ scale.

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Bezaleel
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[*] posted on 5-11-2020 at 03:38


Quote: Originally posted by Tsjerk  
Adding methanol to a potassium hydrogen sulfate solution works and is nearly quantitative.

http://www.sciencemadness.org/talk/viewthread.php?tid=79548&...

That is of course a great find. But will KHSO4 remain available OTC when H2SO4 won't?
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njl
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[*] posted on 5-11-2020 at 04:32


Probably, considering it has commercial use as a fertilizer, pool chemical, food additive etc. Those tend to be difficult to regulate.
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